Ionic and Metallic Bonding
Kevin Pyatt, Ph.D.
Donald Calbreath, Ph.D.
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AUTHORS
Kevin Pyatt, Ph.D.
Donald Calbreath, Ph.D.
EDITORS
Donald Calbreath, Ph.D.
Max Helix
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Printed: August 20, 2014
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Chapter 1. Ionic and Metallic Bonding
C HAPTER
1
Ionic and Metallic Bonding
C HAPTER O UTLINE
1.1
Ions
1.2
Ionic Bonds and Ionic Compounds
1.3
Metals and Metallic Bonds
1.4
References
The image above shows the largest gold nugget ever discovered in California, weighing 156 ounces. Gold is widely
used for money, decorative purposes, and various practical applications in fields such as dentistry, electronics, and
medicine. Its high malleability, ductility, ability to conduct electricity, and resistance to corrosion and most other
chemical reactions make it a highly desirable material in things like electric wiring, colored-glass production, and
corrosion-resistant jewelry and dishes. Gold is also one of the few metals that occurs naturally in its pure form.
Due to their tendency to form cations, most naturally occurring metals are found as part of ionic compounds. For
example, aluminum is the most abundant metal on earth, but it is rarely found in its elemental form. Instead, it is
found as the mineral bauxite, an ionic substance composed of aluminum cations and oxygen anions. Pure aluminum
must be extracted from minerals like bauxite through chemical means. Pure metals have very different properties
than the ionic compounds that they can form with various nonmetals. Additionally, metals can be mixed together to
make alloys that have different properties than either parent metal. In this chapter, we will investigate and compare
some of these different types of substances.
Chris Ral ph (User:Reno Chris/Wikipedia). commons.wikimedia.org/wiki/File:Stringer156_nugget. j pg. Public Domain.
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1.1. Ions
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1.1 Ions
Lesson Objectives
•
•
•
•
Explain how the periodic table can be used to predict the likely charges for ions of a given element.
Depict atoms and ions using electron dot notation.
Describe the octet rule and how it is used to explain chemical behavior.
Define and describe the arrangement of the valence electrons for a given chemical species.
Lesson Vocabulary
• octet rule: States that elements tend to form compounds in ways that give each atom eight valence electrons.
• Lewis electron dot structure: A diagram for a chemical substance in which each element is represented by
its symbol and each valence electron is represented by a single dot.
• isoelectronic: Two atoms or ions with the same number of electrons.
• cation: A positively charged ion.
• anion: A negatively charged ion.
Check Your Understanding
• How do ions differ from atoms? What types of elements form cations, and what types of elements form
anions?
Introduction
As we studied in our chapter on the periodic table, we saw that elements share a number of important properties
with other elements found in the same group. The chemical behavior of a given element is largely dictated by the
configuration of its valence electrons. Many elements have a tendency to gain or lose electrons in order to achieve a
more stable configuration. When a neutral atom gains or loses electrons, it becomes an ion. In this lesson, we will
look at ways to predict what type of ion a given element is likely to form.
Octet Rule
The noble gases are unreactive because of their electron configurations. American chemist Gilbert Lewis (18751946) used this observation to explain the types of ions and molecules that are formed by other elements. He called
his explanation the octet rule. The octet rule states that elements tend to form compounds in ways that give each
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Chapter 1. Ionic and Metallic Bonding
atom eight valence electrons. An exception to this rule is the elements in the first period, which are particularly
stable when they have two valence electrons. A broader statement that encompasses both the octet rule and this
exception is that atoms react in order to achieve the same valence electron configuration as that of the nearest noble
gas. Most noble gases have eight valence electrons, but because the first principal energy level can hold a maximum
of two electrons, the first noble gas (helium) needs only two valence electrons to fill its outermost energy level. As a
result, the nearby elements hydrogen, lithium, and beryllium tend to form stable compounds by achieving a total of
two valence electrons.
There are two ways in which atoms can satisfy the octet rule. One way is by sharing their valence electrons with
other atoms, which will be covered in the next chapter. The second way is by transferring valence electrons from
one atom to another. Atoms of metallic elements tend to lose all of their valence electrons, which leaves them with
an octet from the next lowest principal energy level. Atoms of nonmetallic elements tend to gain electrons in order
to fill their outermost principal energy level with an octet.
Electron Dot Diagrams
A common way to keep track of valence electrons is with Lewis electron dot structures. In an electron dot structure,
each atom is represented by its chemical symbol, and each valence electron is represented by a single dot. Note that
only valence electrons are shown explicitly in these diagrams. For the main group elements, the number of valence
electrons for a neutral atom can be determined by looking at which group the element belongs to. In the s block,
Group 1 elements have one valence electron, while Group 2 elements have two valence electrons. In the p block,
the number of valence electrons is equal to the group number minus ten. Group 13 elements have three valence
electrons, Group 14 elements have four, and so on. The noble gases in Group 18 have eight valence electrons, and
the full outer s and p sublevels are what give these elements their special stability. Representative dot diagrams are
shown in the Figure 1.1:
FIGURE 1.1
The image shown here displays dots circling each elemental symbol. Elements
will typically gain, lose or share electrons
to achieve an octet. Only one group of elements (the noble gases) has a complete
octet as neutral atoms.
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1.1. Ions
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Cations and Anions
Metals will typically lose electrons to achieve stability, while non-metals typically gain electrons to achieve stability.
Two atoms or ions with the same number of electrons are referred to as isoelectronic.
FIGURE 1.2
Cations and Anions
Cations
A positively charged ion is called a cation. Main group metals will typically form ions by losing enough electrons
to become isoelectronic with the nearest noble gas. For example, lithium, whose configuration is [1s2 2s1 ], will
typically lose one electron to become isoelectronic with helium, which has a configuration of [1s2 ] (see Figure 1.1).
Li → Li+ + e−
[He]2s1
[He]
Similarly, beryllium has 4 electrons (with the configuration [1s2 2s2 ]), so it prefers to lose two electrons, in order to
become isoelectronic with helium (again, [1s2 ]).
Be → Be2+ + 2e−
[He]2s2
[He]
Transition Metal Cations
As we saw in our chapter on the periodic table, the valence electrons for transition metals are variable, and electrons
in the highest occupied d orbitals (which are not part of the valence shell) may or may not be lost in the formation of a
transition metal cation. As a result, many transition metals commonly form more than one type of cation, depending
on how many d electrons are lost. Figure 1.3 depicts some of the typical electron arrangements for the transition
elements.
Anions
A negatively charged ion is called an anion. Nonmetals will typically form ions by gaining enough electrons to
become isoelectronic with the nearest noble gas. For example, fluorine has 7 valence electrons and is one electron
away from being isoelectronic with neon, which has a stable noble gas electron configuration (see Figure 1.1).
F
[He]2s2 2p5
+ e− →
F−
[He]2s2 2p6 or [Ne]
Oxygen has 6 valence electrons in its ground state. Remember that ground state refers to the neutral atom in which
the electrons occupy the lowest possible energy positions. Oxygen is two electrons away from being isoelectronic
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Chapter 1. Ionic and Metallic Bonding
FIGURE 1.3
This image shows the arrangement of electrons in their native, chemically
neutral state. Notice that chromium and manganese have a half-filled d
shell. Copper and zinc have fully filled d shells.
with the nearest noble gas. Oxygen will therefore form ions by gaining two electrons to become isoelectronic with
neon, as shown below:
O
[He]2s2 2p4
+ 2e− →
O2−
[He]2s2 2p6 or [Ne]
Similarly, nitrogen has five valence electrons in it ground state, which is three electrons away from the nearest noble
gas. Nitrogen can gain three electrons to become isoelectronic with neon:
N
[He]2s2 2p3
+ 3e− →
N 3−
[He]2s2 2p6 or [Ne]
Example 8.1
Write the ground state configuration for the nonmetal sulfur, and predict the ion it must form to be isoelectronic with
the nearest noble gas.
Answer:
The ground state configuration for the nonmetal sulfur is written as: 1s2 2s2 2p6 3s2 3p4 . Sulfur has 16 electrons. The
nearest noble gas to sulfur is argon, which has an electron configuration of: 1s2 2s2 2p6 3s2 3p6 . To be isoelectronic
with argon, which has 18 electrons, sulfur must gain two electrons. Therefore sulfur will form a 2- ion, becoming
S2− .
Lesson Summary
• Atoms or groups of atoms that carry an overall electrical charge are referred to as ions. Cations can be formed
when a neutral species loses electrons, while anions are formed when a neutral species gains electrons.
• Particularly for main group elements, the number of electrons a given element has in its outer (valence) shell
largely determines the chemical behavior of that element.
• The octet rule states that atoms will lose, gain, or share electrons to achieve the electron configuration of the
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1.1. Ions
•
•
•
•
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nearest noble gas (8 valence electrons, except for helium, which has 2).
Electron dot diagrams are used to help us visualize the arrangement of valence electrons in a given chemical
species.
When an element loses one or more electrons, a cation is formed. Metals typically become cations when they
interact with other chemical species.
Some transition metals can produce ions with multiple different charges due to the optional participation of d
electrons.
When an element gains one or more electrons, an anion is formed. Nonmetals typically become anions when
they interact with other chemical species.
Lesson Review Questions
1. Draw electron dot diagrams for one metal and one nonmetal.
2. Predict whether each of the following is more likely to become a cation or an anion.
(a)
(b)
(c)
(d)
(e)
Ca
Na
F
Br
S
3. Write the ground state electron configurations for the following elements, and predict the ion that will form
when each atom becomes isoelectronic with the nearest noble gas.
(a)
(b)
(c)
(d)
Be
Mg
O
Al
4. Describe the change that is happening when Li → Li+ .
5. What would be the electron configuration for Mg− ? Use the octet rule to explain why this is not likely to be a
very stable ion.
6. What would be the electron configuration for F+ ? Use the octet rule to explain why this is not likely to be a
very stable ion.
7. Element X has a total of 16 electrons.
(a) Write the electron configuration for element X.
(b) How many electrons away from a complete octet is this element?
(c) Make a prediction about the ion that this element might form in an ionic compound.
8. Element Z has a total of 12 electrons.
(a) Write the electron configuration for element Z.
(b) How many electrons away from a complete octet is this element?
(c) Make a prediction about the ion that this element might form in an ionic compound.
9. Write the ground state configuration for the following elements. Then, show how the element ionizes to
become isoelectronic with the nearest noble gas. Example: Mg [1s2 2s2 2p6 3s2 ]
Mg → Mg2+ + 2e−
[Ne]2s2
(a) Na
(b) Ca
(c) N
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[Ne]
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Chapter 1. Ionic and Metallic Bonding
(d) Br
(e) Al
(f) Se
Further Reading / Supplemental Links
• Pough, Frederick. 1988. Rocks and Minerals, Peterson Field Guides. Boston: Houghton Mifflin.
Points to Consider
• So far, we have been discussing the fact that cations and anions form when electrons are lost or gained,
respectively. However, the electrons must be lost to something or gained from something. Where are electrons
lost to, and where do they originate from?
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1.2. Ionic Bonds and Ionic Compounds
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1.2 Ionic Bonds and Ionic Compounds
Lesson Objectives
• Describe the general properties that distinguish ionic compounds from other substances.
• Define and give examples of ionic compounds. Be able to predict which elements are likely to form ionic
compounds with each other.
• Describe the crystal lattice structures adopted by ionic compounds.
• Define lattice energy and explain what it measures.
Lesson Vocabulary
• ionic bond: The resulting attraction between the positively charged cations and negatively charged anions
• crystal lattice: A three-dimensional sturcture formed by ions in order to maximize the number of attractive
interactions while minimizing the repulsive ones.
• lattice energy: The amount of energy needed to completely pull apart an ionic substance into isolated ions.
• dissolution: Occurs when water interacts with the ions in the crystal lattice, causing the lattice to break apart.
Check Your Understanding
1. Give some examples of commonly encountered ions.
2. Draw electron dot diagrams for atoms of the following elements:
a. calcium
b. oxygen
Introduction
In the last section, we saw that elements may lose or gain electrons to become isoelectronic with the nearest noble
gas. Where do the electrons go when an element loses them to become a cation? Where do electrons come from
when an element gains them to become an anion? For an atom to gain or lose electrons, there must be an interaction
between two different chemical species. If electrons are fully exchanged, then we consider this interaction to be
ionic. The resulting attractions between the positively charged cations and the negatively charged anions are referred
to as ionic bonds.
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Chapter 1. Ionic and Metallic Bonding
Ionic Bonds
As we saw in earlier chapters, the electrons in the outermost (valence) shell of an atom are largely responsible for the
ways in which that atom will interact with other elements. For example, Figure 1.4 shows the electron configurations
of sodium (11 e− , 1 valence e− ), neon (10 e− , 8 valence e− ), and fluorine (9 e− , 7 valence e− ).
FIGURE 1.4
This image shows the arrangement of
electrons in their ground states for
sodium, neon, and fluorine.
Our model of ionic bonding and chemical reactivity states that sodium and fluorine have a strong driving force to
become isoelectronic with the nearest noble gas, neon. Because sodium needs to lose one electron and fluorine needs
to gain one for this to occur, one atom of sodium can give up its valence electron to a fluorine atom, resulting in
two ions with noble gas configurations matching that of neon ( Figure 1.5). The positive and negative ion are held
tightly together by electrostatic forces, which are strong forces between oppositely charged particles. When large
groups of sodium and fluorine atoms react in this way, the result is the ionic compound, sodium fluoride.
FIGURE 1.5
Electron arrangements for sodium fluoride.
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1.2. Ionic Bonds and Ionic Compounds
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Crystal Lattices
Any ionic compound is composed of extremely large numbers of cations and anions. Each cation is attracted by
all of the anions but repelled by all the other cations, and vice versa. In order to maximize the number of attractive
interactions while minimizing the repulsive ones, the ions form a three-dimensional structure known as a crystal
lattice. There are a variety of lattice forms that ionic compounds can exhibit, but all of them involve a regular,
repeating pattern in which cations and anions are held rigidly in place by various neighboring ions. For example,
sodium fluoride takes the form of a cubic lattice, shown here ( Figure 1.6):
FIGURE 1.6
Crystal Lattice for Sodium Fluoride
Some properties of the crystal form that are exhibited at the atomic level can also be seen at the macroscopic level.
Due to the cubic arrangement of ions in sodium fluoride, a single pure crystal of this compound will tend to have
smooth faces at right angles to one another ( Figure 1.7).
FIGURE 1.7
Crystals of Villiaumite, a rare mineral
composed of sodium fluoride.
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Chapter 1. Ionic and Metallic Bonding
Lattice Energy
There are a number of different ways to measure the strength of a given crystal lattice. One way would be to measure
the amount of energy needed to completely pull apart an ionic substance into isolated ions. This value, known as
the lattice energy, cannot be measured directly, but it can be calculated based on measured energy changes for
other more feasible processes. The lattice energy of an ionic solid provides us with one way to measure the relative
strength of the ionic bonds in that compound. Table 1.1 shows the lattice energies for various ionic substances:
TABLE 1.1: Lattice Energies for Some Ionic Compounds
Compound
LiF
LiCl
LiI
NaF
NaCl
NaBr
NaI
KF
KCl
Lattice Energy (kJ/mol)
1030
834
730
910
788
732
682
808
701
Compound
KBr
CsCl
CsI
MgCl2
SrCl2
MgO
CaO
SrO
ScN
Lattice Energy (kJ/mol)
671
657
600
2326
2127
3795
3414
3217
7547
Properties of Ionic Compounds
Ionic compounds exhibit certain properties, some of which are listed below:
•
•
•
•
All ionic compounds form crystals.
Ionic compounds tend to have high melting points and boiling points.
Ionic compounds are very hard and very brittle.
Ionic compounds conduct electricity when dissolved in water.
The last property above requires some additional explanation. We are all familiar with the process of dissolution
on a large scale. If you stir a spoonful of salt into a glass of water, the salt crystals are broken down and seem to
disappear into the water. On the atomic level, the dissolution of an ionic compound occurs when water interacts
with the ions in the crystal lattice, causing the lattice to break apart ( Figure 1.8):
Once the ions are dissolved, the presence of charged particles distributed throughout the liquid allows the solution
to conduct electricity ( Figure 1.9). The more ions that are freed from the lattice, the more conductive the solution
will be.
Below is a summary of some common ionic compounds and their practical applications ( Table 1.2):
TABLE 1.2: Common Examples of Ionic compounds
Formula
NaCl
NaHCO3
NaOH
NaF
Name
Sodium chloride
Sodium hydrogen carbonate
Sodium hydroxide
Sodium fluoride
Common name
Table salt
Baking soda
Uses
Food additive
Mild cleaner, antacid
Lye
n/a
Drain cleaner
Active ingredient
toothpaste
in
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1.2. Ionic Bonds and Ionic Compounds
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TABLE 1.2: (continued)
Formula
NaOCl
12
Name
Sodium hypochlorite
Common name
Bleach
Uses
Mild or strong cleaner,
disinfectant
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Chapter 1. Ionic and Metallic Bonding
FIGURE 1.8
Dissolution of NaCl.
FIGURE 1.9
Conductivity of ionic solutions.
Lesson Summary
• Ionic bonds are electrostatic attractions between two oppositely charged ions. Ions can be formed and then
bonded when metal atoms donate their valence electrons to nonmetal atoms.
• The ions in ionic compounds are arranged in rigid three-dimensional patterns called crystal lattices. The
crystal lattice that is formed is a characteristic property of a given compound.
• We can indirectly measure the energy necessary to break apart a given lattice into its isolated ions. We call this
value the lattice energy, and it gives us one way to measure the strength of the ionic bonds in that compound.
• Ionic compounds have the following properties: (1) they form crystals; (2) they have high melting/boiling
points; (3) they are hard and brittle; (4) they can conduct electricity when dissolved in water.
• Dissolution is a process in which water interacts with the ions in a crystal lattice, causing the lattice to break
apart.
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1.2. Ionic Bonds and Ionic Compounds
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Lesson Review Questions
1. How do the electrons from two atoms interact in an ionic bond?
2. Predict the formulas for the ionic compounds formed when each of the metals in the Table 1.3 reacts with
each nonmetal.
TABLE 1.3: Reactions
Oxygen
Sulfur
Chlorine
Calcium
Sodium
Aluminum
3. What is dissolution?
4. If sodium chloride is placed in water, it will completely dissociate into its ionic components, described by
−
the dissociation equation: NaCl(s) → Na+
(aq) + Cl(aq) . (The (aq) signifies an aqueous solution, or a solution in
which the ions are dissolved in water.) Write a similar dissociation equation for the solid ionic compound
calcium chloride.
5. Which physical properties of ionic compounds can be attributed to the crystal lattice structure?
6. How does lattice energy relate to the strength of an ionic compound?
7. True or false: The high melting points of ionic solids suggest that ionic bonds are fairly weak.
8. Using Table 1.1 as a reference, what trend can be recognized between lattice energy and the characteristics of
the ions which comprise the compound? For example, NaF, NaCl, NaBr, NaI have lattice energies (kJ/mol)
of 910, 788, 732, and 682, respectively. What is different between the anions that may be causing such
differences?
Further Reading / Supplemental Links
• Animation of Sodium chloride dissolution: http://www.mhhe.com/physsci/chemistry/essentialchemistry/fla
sh/molvie1.swf
Points to Consider
• We have generally assumed that ionic compounds are composed of metal cations and nonmetal anions. While
this is common, there are ionic compounds in which no metal is involved. For example, the ammonium cation
is positively charged but does not involve any metal atoms. Similarly, some polyatomic anions are not solely
comprised of nonmetallic atoms. What are some examples of these anions?
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Chapter 1. Ionic and Metallic Bonding
1.3 Metals and Metallic Bonds
Lesson Objectives
•
•
•
•
Describe the general properties of metals compared to other element types.
Describe the arrangement of atoms in metallic substances.
Describe the behavior of electrons in metals.
Define and give examples of alloys.
Lesson Vocabulary
•
•
•
•
•
•
•
•
malleable: When pure metals are able to be stamped, pressed, or rolled into thin sheets.
ductile: Metal that can be stretched, bent, or twisted without breaking.
toughness: The ability of a material to withstand shock and to be deformed without rupturing.
luster: When pure metals tend to be shiny in appearance.
corrosion: The gradual degradation of a material due to its exposure to the environment.
metallic bond: The attraction of the stationary metal cations to the surrounding mobile electrons.
alloy: A mixture of pure metals.
amalgam: An alloy that is mostly composed of mercury.
Check Your Understanding
1. Identify the ions that make up the following compounds:
a. NaCl
b. BaSO4
c. K2 O
2. How many valence electrons do the neutral atoms of metals in Groups 1, 2, and 3 in the periodic table have?
Introduction
Metals represent approximately 25% of the elemental makeup of the Earth’s crust. The bulk of these metals,
primarily aluminum, iron, calcium, sodium, potassium, and magnesium, are typically found in combined form.
The most abundant metal is aluminum, which occurs almost exclusively as the ionic mineral bauxite. The other
most common metals, including iron, sodium, potassium, magnesium, and calcium, are also found primarily as the
cationic portion of an ionic compound. Very few metals actually occur naturally as pure substances. The ones that
do are often referred to as precious or semi-precious metals.
As pure substances, metals are tough, yet malleable. They are strong, and some of them are quite resistant to
corrosion. They are also good conductors of electricity and heat. Due to these and other useful properties, pure
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1.3. Metals and Metallic Bonds
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metals have been valued for millennia. In this lesson, we are going to investigate a few properties of metals and the
chemical reasons behind some of these characteristics.
Properties of Metals
Physical Properties
Most pure metals share a number of physical properties. Metals are good conductors of electricity and heat. They are
also malleable, which means that they can be stamped, pressed, or rolled into thin sheets. For example, aluminum
foil can be made into sheets that are only 13 µm thick, and gold (the most malleable pure metal) can be hammered so
thin that it is practically transparent. Metals also tend to be ductile, which means that they can be stretched, bent, or
twisted without breaking. The copper wire shown in Figure 1.10 is an example of this. Both of these properties are
facets of toughness, which is a term that describes the ability of a material to withstand shock and to be deformed
without rupturing.
FIGURE 1.10
This image shows a variety of different
copper wires.
Copper is a commonly
used substance for wire because it is
highly conductive and ductile but also very
abundant (and therefore inexpensive).
Pure metals tend to be shiny in appearance; this property is referred to as luster. Due to our everyday experiences,
we may think of metals as being mostly dull gray in color. However, this is due not to the pure metal but to a surface
layer in which the pure metal has formed an ionic compound, usually with oxygen atoms from either air or water.
Most pure metals are silver-white, but some of the heavier ones (most notably, gold) take on a yellowish hue.
Chemical Properties
We have already discussed some of the chemical properties of pure metals. They have just a few valence electrons
(generally 1-3), which tend to be fairly easy to remove due to metals low ionization energy and electronegativity
values. As a result, they frequently form ionic compounds by transferring their valence electrons to nonmetallic
atoms, which use these extra electrons to complete their valence shells and achieve noble gas configurations.
The driving force to combine with nonmetals to create ionic compounds varies quite a bit between different metals.
Some pure metals, like cesium and potassium, are so eager to react that they must be stored under oil to avoid an
immediate reaction with the oxygen present in air. Others, like platinum and gold, are stable enough that they can
be found in nature as pure metals rather than as the cationic portion of an ionic compound. Gradual degradation of a
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Chapter 1. Ionic and Metallic Bonding
material due to its exposure to the environment is known as corrosion. Metals like gold and platinum are unusually
resistant to corrosion, which makes them especially valuable for both structural and decorative purposes.
Metals have a wide range of melting points, but most are quite high. Only one metal (mercury) melts below room
temperature. Others (such as gallium) are solid at room temperature but would melt at body temperature, so they
can be melted simply by holding them in your hand. On the other end of the spectrum, tungsten has a melting point
of 3422°C. Figure 1.11 shows the melting points of various elements in their most common pure form.
FIGURE 1.11
Melting Points of the Metallic Elements
The "Sea of Electrons"
The reason metals behave the way they do can largely be explained by the ways that metal atoms bond together
to make a solid material. Pure metals are crystalline solids, but unlike ionic compounds, every point in the crystal
lattice is occupied by an identical atom. The electrons in the outer energy levels of a metal are mobile and capable of
drifting from one metal atom to another. This means that the metal is more properly viewed as an array of positive
ions surrounded by a "sea" of mobile valence electrons ( Figure 1.12). Electrons that are capable of moving freely
throughout the empty valence orbitals of the metallic crystal are said to be delocalized. A metallic bond is the
attraction of the stationary metal cations to the surrounding mobile electrons.
FIGURE 1.12
Electron Sea Illustration
This model for metallic bonding explains some of the physical properties of metals. Metals conduct electricity and
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1.3. Metals and Metallic Bonds
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heat very well because of their free-flowing electrons. As electrons enter one end of a piece of metal, an equal
number of electrons flow outward from the other end, allowing an electrical current to pass through the material
with minimal resistance. Additionally, because the electron "glue" that holds the metal atoms together is very easy
to deform and reshape, bulk metals can be easily hammered, bent, and pulled without breaking apart.
Types of Metals
Precious Metals
A number of relatively rare metals are quite resistant to corrosion. These metals are sometimes referred to as precious
metals due to their scarcity and their ability to remain pure over time. The exact list varies, but metals that are usually
classified as precious include gold, silver, ruthenium, rhodium, palladium, osmium, iridium, and platinum. Some
precious metals are shown in ( Figure 1.13). Compared to other metals, precious metals tend to have relatively high
ionization energies and electronegativity values.
FIGURE 1.13
Precious metals
Rare Earth Metals
The rare earth metals are a set of seventeen chemical elements (the lanthanide series plus scandium and yttrium)
that have particular importance for a variety of industrial processes and are used frequently in modern technology.
Despite their name, rare earth metals are actually relatively abundant in the earth’s crust. However, the extraction of
many of these metals is quite difficult and has made their supply somewhat limited. They are highly sought after for
this reason. Figure 1.14 shows the rare earth metals.
Alloys
In addition to being used in their pure elemental forms, metals can be melted down and combined with other metals
(and sometimes small amounts of nonmetals) to form mixtures known as alloys. The properties of alloys are often
quite different than the properties of the base elements from which they formed. For example, iron is often mixed
with small amounts of carbon or other metals to create steels. By modifying the relative amounts of the added
components, properties like hardness, flexibility, and corrosion resistance can be fine-tuned so that the material is
suitable for a particular application. For example, elemental iron corrodes readily in air and water (see Figure 1.15),
but stainless steel (which is still mostly iron, but contains about 10-12% chromium by mass) resists corrosion to a
large extent. It is used as an exterior building material for extravagant buildings such as that shown in Figure 1.16.
Alloys that are mostly composed of mercury are known as amalgams. Amalgams often have special properties that
stem from the fact that mercury exists as a liquid at room temperature. As a result, metal amalgams are used for a
variety of purposes, including dentistry and the extraction of other pure metals such as gold.
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Chapter 1. Ionic and Metallic Bonding
FIGURE 1.14
Rare earth metals
Lesson Summary
• Physical properties that are common to metals include malleability, ductility, toughness, and luster.
• Chemical properties of metals include the abilities to conduct heat and electricity.
• Many of the properties of metals are due to the presence of metallic bonds, in which metal atoms are held
together by a mutually shared "sea" of valence electrons.
• Metal atoms lose electrons readily to other chemical species, forming cations that can participate in ionic
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1.3. Metals and Metallic Bonds
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FIGURE 1.15
Corroded iron pipe.
FIGURE 1.16
Most of the exterior of the Walt Disney
Concert Hall pictured here is stainless
steel.
bonds.
• Precious metals are relatively rare and often can be found naturally as a pure metal rather than as part of an
ionic compound.
• Rare earth metals are an important set of metals because of their use in a variety of modern industrial
applications.
• Alloys are solid mixtures of two or more metals (sometimes with small amounts of a nonmetal, such as
carbon). The properties of alloys may be quite different than the properties of the pure elements from which
they are formed.
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Chapter 1. Ionic and Metallic Bonding
Lesson Review Questions
1. Define the following properties of metals
(a)
(b)
(c)
(d)
2.
3.
4.
5.
6.
7.
8.
9.
malleability
ductility
toughness
luster
In metal solids, the ______ electrons form a shared sea of electrons.
What is corrosion as it applies to metals?
In general, what can be said of the melting points of metals?
Define a metallic bond.
What is the relationship between the electron arrangement in metals and metals’ physical properties?
What atomic properties distinguish "precious metals" from metals in general?
Why are "rare metals" so valuable?
How are alloys formed?
Further Reading / Supplemental Links
• Metallic Bonding Animation/Video: http://www.youtube.com/watch?v=c4udBSZfLHY
• Use of Different Types of Alloys: http://www.articlesnatch.com/Article/Uses-Of-Different-Types-Of-Allo
ys/599802
Points to Consider
• The following image is of a sword constructed of an alloy known as Damascus steel.
Notice the mottling pattern, reminiscent of flowing water. Damascus steel blades were a product of medieval times
and were highly regarded for their toughness, resistance to shattering, and their capability to be honed and sharpened
with a resilient edge. Even to this day, it is still debated as to exactly how Damascus steel was made.
21
1.4. References
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1.4 References
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
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Christopher Auyeung and Jodi So. CK-12 Foundation . CC BY-NC 3.0
Zachary Wilson. CK-12 Foundation . CC BY-NC 3.0
Jodi So. CK-12 Foundation . CC BY-NC 3.0
Jodi So. CK-12 Foundation . CC BY-NC 3.0
Jodi So. CK-12 Foundation . CC BY-NC 3.0
Ben Mills (User:Benjah-bmm27/Wikimedia Commons). http://commons.wikimedia.org/wiki/File:Sodium-fl
uoride-3D-ionic.png . Public Domain
User:Stickpen/Wikimedia Commons. http://commons.wikimedia.org/wiki/File:Villiaumite-russia.jpg . Public Domain
Christopher Auyeung. CK-12 Foundation . CC BY-NC 3.0
Christopher Auyeung. CK-12 Foundation . CC BY-NC 3.0
Image copyright Flegere, 2014. http://www.shutterstock.com . Used under license from Shutterstock.com
Jodi So. CK-12 Foundation . CC BY-NC 3.0
Christopher Auyeung. CK-12 Foundation . CC BY-NC 3.0
Hi-Res Images of Chemical Elements. http://images-of-elements.com . CC BY 3.0
Hi-Res Images of Chemical Elements. http://images-of-elements.com . CC BY 3.0
Mark (Flickr:RelentlesslyOptimistic). http://www.flickr.com/photos/relentlesslyoptimistic/5032289/ . CC
BY 2.0
User:Arturoramos/Wikimedia Commons. http://commons.wikimedia.org/wiki/File:Walt_Disney_Concert_H
all_Across_Grand.jpg . CC BY 3.0